The document defines key terms related to thermodynamics and calorimetry, including enthalpy, enthalpy of reaction, calorimetry, heat capacity, bomb calorimeter, Hess's law, enthalpy of formation, and standard enthalpy of reaction. It then provides examples of using these concepts to calculate enthalpy changes, heat transferred, and moles or mass of substances involved in various chemical reactions.

1. KEY
GENERAL CHEMISTRY-I (1411)
S.I. # 21
1. Define the following:
a. Enthalpy: (H) H = E + PV where P is pressure and V is volume, and H is the
heat gained or lost by the system at constant pressure ∆H = qp. For endothermic reactions
∆H > 0, for exothermic reactions ∆H < 0.
b. Enthalpy of reaction: ∆Hrxn = H(products) – H(reactants). This is a chemical
characteristic of the reaction. Two rules: #1: ∆H is proportional to the amount of reactant
that reacts. #2: Reversing a reaction changes the sign of ∆H.
c. Calorimetry: the experimentally measured amount of heat transferred between
the system and the surroundings.
d. Heat capacity: the amount of heat required to raise the temperature by 1 K.
The heat capacity for one mole (1mol) of pure substance is called its molar heat capacity.
For one gram (1g) of the substance, we use the term specific heat (s). The heat absorbed
by a system can be found by q = (s) (m) (∆T), where m is the mass.
e. Bomb calorimeter: used to measure the heat evolved in combustion reactions
at constant volume. Here q = ∆E.
f. *Hess’s Law: states that if a reaction is carried out in a series of steps, ∆H for
the reaction will be equal to the sum of the enthalpy changes for the steps.
g. Enthalpy of formation: (∆Hf) of a substance is the enthalpy change for the
reaction in which the substance is formed from its constituent elements.
h. Standard Enthalpy change: (∆H°) is the enthalpy change when all reactants
and products are at 1 atm pressure and a specific temperature, usually 298 K (25°C).
i. Standard Enthalpy of formation: (∆H°f) of a substance is the change in
enthalpy for the reaction that forms one mole of the substance from its elements in their
most stable form with all reactants and products at 1 atm pressure and usually 298 K. For
any element in its stable state (∆H°f) = 0.
∆H°rxn = Σ n ∆H°f (products) - Σ m ∆H°f (reactants)
2. The decomposition of zinc carbonate, ZnCO3 (s), into zinc oxide, ZnO (s), and
CO2(g) at constant pressure requires the addition of 71.5 kJ of heat per mole of
ZnCO3. a) Write a balanced thermochemical equation for the reaction. b) Draw an
enthalpy diagram fro the reaction.
a. ZnCO3 (s)  ZnO (s) + CO2 (g) ∆H = 71.5 kJ
b. ZnO (s) + CO2 (g)
∆H= +71.5 kJ
ZnCO3 (s)

2. KEY
3. Consider the reaction: CH3OH (g)  CO (g) + 2H2 (g) ∆H = +90.7 kJ
a) Is the heat absorbed or evolved in the course of this reaction? Is the reaction endo
or exo? b) Calculate the amount of heat transferred when 45.0 g of CH3OH (g) is
decomposed by this reaction at constant pressure. c) For a given sample of CH3OH,
the enthalpy change on reaction is 18.5 kJ. How many grams of hydrogen gas are
produced? d) What is the value of ∆H for the reverse of the previous reaction? How
many kilojoules of heat are released when 27.0 g of CO (g) reacts completely with H2
(g) to form CH3OH (g) at constant pressure?
a. The reaction is endothermic so heat is absorbed by the system.
b. (45.0g CH3OH) (1 mol CH3OH / 32.04 g CH3OH)(90.7 kJ / 1mol CH3OH) = 127 kJ
heat transferred (absorbed).
c. (18.5 kJ)(2mol H2 / 90.7 kJ) (2.016 g H2 / 1 mol H2) = 0.822 g H2 produced
d. The sign of ∆H is reversed for the reversed reaction: ∆H = -90.7 kJ
(27.0 g CO)(1mol CO / 28.01gCO)(-90.7 kJ / 1mol CO) = -87.4 kJ heat transferred
(released).
4. Use: 2KClO3 (s)  2 KCl (s) + 3 O2(g) ∆H = -89.4 kJ
For this reaction, calculate ∆H for the formation of a) 0.855 mol of O2 and b) 10.75 g
of KCl.
a) (0.855 mol O2) (-89.4 kJ / 3 mol O2) = -25.5 kJ
b) (10.75 g KCl) (1mol KCl / 74.55 g KCl) (-89.4 kJ / 2mol KCl) = -6.45 kJ
5. When 418 J of heat is added to a gas under constant atmospheric pressure, it
expands and does 107 J of work on the surroundings. What are the values of ∆H
and ∆E for this process?
The gas is the system. If 418 J of heat is added, q = +418 J. Work done by the system
decreases the overall energy of the system, so w = -107 J. Therefore,
∆E = q + w = 418 J – 107 J = 311 J
∆H = q = 418 J (at constant pressure)
6. The specific heat of ethylene glycol is 2.42 J/g-K. How many J of heat are needed
to raise the temperature of 62.0 g of ethylene glycol from 15.2 °C to 40.8 °C?
(62.0 g ethylene glycol) (2.42 J / g K) (40.8°C – 15.2°C) = 3.84x103 J